Salt production methods seek to maximize the various commercial products that can be recovered from salt mixtures, but struggle to find methods that are as efficient and as cost-effective as possible. Moreover, there is growing awareness of the need to develop sustainable processes that are environmentally-friendly.
Sodium chloride can be recovered from a salt water source such as sea water, salt lake water, brine wells, underground deposits, etc., through evaporation of water and manual recovery of the sodium chloride. For example, sodium chloride is generally the first significant compound precipitated from the evaporation of sea water in salt ponds. However, other mixed salts can then be precipitated from the remaining brine. These mixed salts can comprise any number of salts, including sodium chloride (also called Halite or NaCl), magnesium sulfate (MgSO4), magnesium chloride (MgCl), potassium chloride (KCl), magnesium sulfate heptahydrate (Epsom salts or Epsomite, MgSO4.7H2O), magnesium sulfate hexahydrate (Hexahydrate, MgSO4.6H2O), magnesium sulfate monohydrate (Kierserite, MgSO4.1H2O), as well as double salts such as potassium sulfate dimagnesium sulfate (Langbeinite, K2SO4.2MgSO4), potassium chloride plus sodium chloride (Sylvinite, KCl+NaCl), sodium magnesium sulfate (Bloedite or Astrakanite. Na2SO4.MgSO4.4H2O), sodium sulfate (NaSO4) and potassium magnesium chloride sulfate (Kainite, 4KCl.4MgSO4.11H2O). These salts can be either in solid form or can be found in solution as separate ions, such as sodium, magnesium, potassium, calcium, chloride, sulfate, and carbonate.
Several of these compounds have commercial value, but their recovery from these types of salt mixtures can be challenging. Further recovery of sodium chloride has obvious commercial value. Magnesium sulfate salts and hydrates such as Epsom salts also have several uses including as bath salts, laxatives, fertilizers, animal feeds and detergent fillers. Magnesium chloride can be used as a dust suppressant, deicer, anti-icer, or a pre-wetting agent. Some of the issues encountered when trying to recover these products from the salt mixture include the similar chemistry between sodium and magnesium ions, the tendency of Epsomite and sodium sulfate to form double salts such as Astrakanite, the various hydrate forms that magnesium sulfate can be found in, the energy requirements to separate the compounds, and the residual by-products produced from known methods.
Due to these significant limitations, few commercial processes exist for the recovery of such salts from naturally-occurring salt mixtures. As a result, salts such as magnesium sulfate are usually made via other means. The most common method of producing magnesium sulfate is through the chemical process of reacting sulfuric acid with either magnesium oxide or magnesium carbonate. This industrial method is expensive, and may also result in heavy metals in the product due to contamination in the feed stocks, which then have to be removed via further processing.
Another common process used to produce compounds such as magnesium sulfate and Epsom salts from salt mixtures such as salt ponds is to subject the salt mixture to further evaporation, refrigeration, and re-crystallization (see, e.g., Fernandez-Lozano, J. A. “Recovery of Epsomite and Sylvite From Seawater Bittern by Crystallization.” Fourth Symposium on Salt, Alan H. Coogan, The Northern Geological Society, Inc. Cleveland, Ohio, vol. 2, pp. 501-510 (1974)). Following evaporation, the mixture is diluted by adding fresh water to control precipitation of sodium chloride. It is then cooled below ambient temperatures (typically below −5° C.) to separate the magnesium sulfate compounds from the other compounds, followed by re-crystallization. Among the drawbacks of this process are the costs involved in the refrigeration process and the fouling of the heat exchanger due to encrustation of solids on the surface (often referred to as scaling). Another drawback is that the solid-liquid separations are more difficult at cooler temperatures due to increased viscosity of the magnesium chloride brine. In addition, the yield (or product recovery) is significantly less in this refrigeration process than in the present invention. Moreover, this refrigeration process is suitable only for the recovery of Epsom salt and liquid magnesium chloride, whereas the present invention can recover Epsom salt, sodium chloride and liquid magnesium chloride.
One process disclosed in U.S. Pat. No. 5,281,242 involves treating a starting two compound mixture of magnesium sulfate heptahydrate (MgSO4.7H2O) and sodium chloride (NaCl) by applying heat to the mixture in order to convert the magnesium sulfate heptahydrate into clusters of lower hydrate compounds such as magnesium sulfate hexahydrate (MgSO4.6H2O), magnesium sulfate monohydrate (MgSO4.1H2O) or anhydrous magnesium sulfate (MgSO4). Pressure is then applied to the magnesium sulfate clusters to form fine crystals, followed by size separation (such as sieve or screening) and recovery of the magnesium sulfate crystals from the sodium chloride crystals. In order to obtain Epsomite, the magnesium sulfate crystals need to be further processed by re-crystallizing the crystals. One limitation of this process is that it is limited to a starting material consisting only of Epsomite and Halite (many naturally-occurring starting materials consist of several different types of salts as described above). Another limitation is that the temperature and pressure applied to the hydrate clusters and crystals has to be controlled in order to obtain the appropriate magnesium sulfate hydrate salt. Finally, this is a solid-solid separation process that does not involve the use of liquids in contrast to the present invention.